chemistry unit 2

Isotopes

isotopes ­ atoms of the same element with different mass numbers

to have different mass number, the atoms must have different numbers

of neutrons

neutrons don't impact the reactivity of the element so they occupy the

same place in the periodic table

Atomic Absorption and Emission Line Spectra

When EM radiation is passed through a collection of atoms some of the

radiation is absorbed and used to excite the atoms from a lower energy

level to a higher energy level.

An absorption spectrum shows the radiation absorbed as atoms move

from lower to higher energy levels.

An emission spectrum is produced when an atom moves from a higher

to a lower level.

Different elements have different line spectra, this information was be

used to identify unknown elements.

Evidence for the Bohr Model

Niels Bohr proposed that an electron moves into an orbit or higher

energy level further from the nucleus when an atom absorbs energy.

The excited state produced is unstable and the electron soon falls

back to the lowest level or ground state.

The energy the electron gives out when it falls into lower levels is in

the form of EM radiation

One packet of energy (quantum) or photon, is released for each

electron transition.

The energy of the photon is proportional to the frequency of the

radiation.

The Hydrogen Spectrum

H atom gives out energy when an electron falls from a higher to a

lower energy level.

H produces visible light when the electron falls to the second energy

level (n=2). The transitions to the first energy level (n=1) correspond

to a higher energy change and are in the UV region of the spectrum.

IR radiation is produced when an electron falls to the 3rd+ E level.

The pattern of lines gives us a picture of the energy levels in the

atom.

The lines converge at higher energies because the energy levels

inside the atoms are closer together at high energy.

When an electron is at the highest energy it's no longer in the atom

and the atom has been ionized.

The energy needed to remove an electron from the ground state of

each atom in a mole of gaseous atoms, ions, or molecules is called

the ionization energy.

The Uncertainty Principle

An electron's trajectory can't be measured, as any attempt to measure

an electron's position will disturb it's motion.

Heisenberg's Uncertainty Principle ­ we can't know where an electron

is at any given moment in time, we can only hope to predict where it is

likely to be

Schrodinger Model of the H Atom

Schrodinger proposed that a wave equation could be used to describe

the behaviour of an electron in the same way that a wave equation

could be used to describe the behaviour of light.

The equation can be applied to multi­electron systems and its

solutions are known as atomic orbitals.

Atomic orbitals ­ a region around an atomic nucleus in which there is a

90% probability of finding the electron. The shape of the orbital will

depend on the energy of the electron. When an electron is in an orbital

of higher energy it will have a higher probability of being found further

from the nucleus.

Atomic Orbitals

2s has the same symmetry as a 1s orbital but extends over a larger

volume so 2s are, on average, further from the nucleus than

electrons in 1s orbitals and are at higher energy.

Electron Spin and the Pauli Exclusion Principle

Pauli Exclusion Principle ­ no more than 2 electrons can occupy any

one orbital, and if two electrons are in the same orbital they must spin

in opposite directions (to overcome their mutual repulsion)

Aufbau Principle: orbital diagrams

Aufbau Principle ­ electrons are placed into orbitals of lowest energy

first.

Hund's Third Rule

Hund's Third Rule ­ if more than 1 orbital in a sub­level is

available, electrons occupy different orbitals with parallel spins

This configuration minimizes the mutual repulsion between the

electrons

The Relative Energy of the Orbitals Depends on the Atomic Number

The energy of an orbital depends on the attractions between the

electrons and the nucleus and inter­electron repulsions.

3d and 4s levels are very close in energy and their relative separation

is very sensitive to inter­electron repulsion.

4s is filled before 3d but electrons are first lost from 4s before 3d.